# collision theory

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### Reaction Rates Prepared by Barbara Andreswith A. NiedzvieckiApril 5, 2004@ Centennial Regional High School :

Reaction Rates Prepared by Barbara Andreswith A. NiedzvieckiApril 5, 2004@ Centennial Regional High School

### Factors which influence reaction rates: :

Factors which influence reaction rates: Temperature Presence of a catalyst Concentration of reactants Surface area of reactants Type of reactants

### How can we explain these factors? :

How can we explain these factors?

Collision Theory

### Reaction rate depends on: :

Reaction rate depends on: # of collision per unit time Success rate

### Success depends on the: :

Success depends on the: Collision geometry (orientation) Energy of the collision

### Collision Geometry :

Collision Geometry

### H2(g) + I2(g) 2HI(g) Side-to-Side Collision :

H2(g) + I2(g) 2HI(g) Side-to-Side Collision

### Slide 9:

H2(g) + I2(g) X 2HI(g) End-to-End Collision

### Slide 10:

H2(g) + I2(g) X 2HI(g) Glancing Collision

### HCl(g) + C2H4(g) products :

HCl(g) + C2H4(g) products For this example: Hydrogen end of H-Cl approaches C=C double bond ….. Success All other collisions …..“rebound”

### Which of the following collisions will be successful? :

Which of the following collisions will be successful?  X X X

### Energy of the Collision :

Energy of the Collision

### Slide 14:

Particles (molecules & atoms) must collide with a minimum energy Minimum energy called the activation energy of the reaction

### Remember, the energy profile for an exothermic reaction looks like this: :

Remember, the energy profile for an exothermic reaction looks like this:

### In order for a reaction to occur the: :

In order for a reaction to occur the: activation energy “barrier” must be overcome collision energy  activation energy bond breaking & bond formation (reaction) collision energy < activation energy “rebound” (no reaction)

### For any reacting system: :

For any reacting system: Particles have a very wide range of energies What proportion of the particles have high enough energies to react when they collide?

### Slide 18:

Maxwell-Boltzmann Distribution (gases) Area under the curve = total number of particles present

### Slide 19:

Maxwell-Boltzmann Distribution (gases)

### To enable low energy particles to react we have to: :

To enable low energy particles to react we have to: change the shape of the curve or move the activation energy to the left

### How can we change the shape of the curve? :

How can we change the shape of the curve? Change the temperature of the reaction

### Slide 22:

How can we overcome the activation energy “barrier”?

### Slide 23:

Effect of Temperature on Reaction Rates

### Temperature  Reaction Rate  :

Temperature  Reaction Rate  For many reactions happening at around room temperature, the rate of reaction doubles for every 10°C rise in temperature.

### Why does this occur? :

Why does this occur? Increased collision frequency Particles only react when they collide. Heat a substance … the particles move faster Collide more frequently Increased rate of reaction

### Slide 26:

Increased kinetic energy Overcome activation energy “barrier” Heat a substance … the particles move faster Collisions are more “intense” Increased rate of reaction.

### Temperature  … graph shape altered :

Temperature  … graph shape altered Area under a curve = count of # of particles “T + t” area doubled Doubled rate of reaction

### Slide 28:

Effect of a Catalyst on Reaction Rates

### Catalysts are substances that: :

Catalysts are substances that: speed up a reaction participate in the reaction steps are chemically unchanged at the end of the reaction have no loss of mass after the reaction

### Catalysts … :

Catalysts … provide an alternative way for the reaction to happen which has a lower activation energy move the activation energy to the left on Maxwell-Boltzmann Distribution graph.

### Maxwell-Boltzmann Distribution :

Maxwell-Boltzmann Distribution

### On an energy profile: :

On an energy profile:

### Slide 33:

Effect of Concentration on Reaction Rates

### Slide 34:

Na2S2O3(aq) + 2 HCl  S(s) + SO2(aq) + 2 NaCl(aq) + H2O(l) Dilute HCl is added to sodium thiosulphate solution …. precipitate of sulphur forms Sodium thiosulphate solution diluted …. precipitate takes longer and longer to form.

### Case 1: Reactions - 2 Particles :

Case 1: Reactions - 2 Particles Must first collide Concentration is higher  the chance of collision are greater.

### Slide 36:

Case 2: Catalyst already working as fast as it can Small amount of a solid catalyst in a reaction High concentration of reactants Catalyst completely “occupied” Adding more reactants has no effect

### Slide 37:

Case 3: Certain Multi-step Reactions

### Slide 38:

Pressure  Concentration  Reaction Rate  Mass of gas … squeezed into a smaller volume Relationship Between Pressure and Concentration

### Can also be explained by the ideal gas law … :

Can also be explained by the ideal gas law …

### Slide 40:

Effect of Surface Area on Reaction Rates

### Surface Area  Reaction Rate  :

Surface Area  Reaction Rate  Particles in the gas or liquid must collide with the solid particles Increasing the surface area of the solid increases the chances of collision taking place

### Slide 42:

2Mg(s)+ 2 HCl(aq)  H2(g)+2 MgCl

### Slide 43:

Effect of Nature of Reactants on Reaction Rates

### We have studied the following two reactions: :

We have studied the following two reactions: C25H52(s) + 38 O2(g)  25 CO2(g) + 26 H2O(g)  2 CH3OH(l) + 3 O2(g)  2 CO2(g) + 4 H2O(g) Based on experimental results, how did the nature of reactants affects the rate of reaction: # of bond broken and formed  reaction rate 

### We have studied the following two reactions: :

We have studied the following two reactions: C25H52(s) + 38 O2(g)  25 CO2(g) + 26 H2O(g)  2 CH3OH(l) + 3 O2(g)  2 CO2(g) + 4 H2O(g) Explain this based on Collision Theory

### We have studied the following two reactions: :

We have studied the following two reactions: C25H52(s) + 38 O2(g)  25 CO2(g) + 26 H2O(g)  2 CH3OH(l) + 3 O2(g)  2 CO2(g) + 4 H2O(g) Explain this based on Collision Theory Greater # of collisions required

The End